AP CHEMISTRY CHAP. 13
HIGHLIGHTS, NOTES AND OBSERVATIONS
Yes, we are familiar with the concept of an ideal gas and the laws that real gases abide by
(at least under atmospheric conditions). Well, it so happens that we can define and discuss
the concept of an ideal solution. To do so, you have to diagnose the intermolecular forces
that are in play as a solution forms. To dissolve a solute you must break the bonds between
the solute particles and to make room within the solvent, you must break at least some of
those bonds. These two processes are inherently endothermic. Possibly balancing this
energy requirement is the exothermic events of solute and solvent particles forming
intermolecular bonds.
This leads to four scenarios (and large possible interpolations): 1) when a solute is dissolved
the energy to break the bonds is slightly more than the energy released as solute-solvent
bond forming occurs. A thermometer measuring this process would record a cooling. (CS2
and (CH3)2CO) 2) When the solute dissolves, the solute-solvent interactions are more
exothermic than the energy required to break the solute bonds and the solvent bonds – the
solution will get slightly warmer. (CHCl3 and (CH3)2CO) 3) A solution forms and the
temperature remains unchanged – what can you infer about the balance of bond breaking
and bond forming? (C6H6 and C7H8) 4) A solute is mixed into a solvent AND NO
SIGNIFICANT DISSOLVING takes place (C6H6 and H2O). What can you infer about the
natures of the solute and the solvent?
You may recall that endothermic reactions are not as commonly apt to occur spontaneously.
You should also recall that another driving force is entropy. The dispersion of the solute, in
the now more disorganized solvent, is a favored process within the context of
thermodynamics and explains spontaneous solution formation even if a net energy input is
required. Throw a little Le Chatelier's principle in at this point: most salts dissolve in water
and are accompanied by a slight cooling (an endothermic process) – therefore, to enhance
dissolving of salts, raise the temperature of water. Make sense???
Scenario 3 above leads to what we call an ideal solution. If you can diagnose molecules for
their intermolecular attraction, you can infer whether an ideal solution can form. Ideal
solutions have predictable properties. Several follow below.
According to Raoult, the ratio of the partial pressure exerted by solvent vapor pressure above
an ideal solution to the vapor pressure of the pure liquid is equal to the mole fraction,  , (do
you recall the definition of mole fraction?) of the solvent. The equation is more commonly
written:
PA   A PA
The benzene-toluene system is an example. Recall the structure of the two chemicals.
Although most solutions behave nonideally (both positive and negative deviations).
Nevertheless, let us focus only on ideal solution behavior.
Practice some mathematical exercises involving the translation of one unit of concentration
to another. These are found on pages 543-547.
COLLIGATIVE PROPERTIES: vapor pressure lowering, boiling point elevation, freezing
point depression and osmotic pressure. These properties are bound together through
their common origin: Each depend only on the number of solute particles present, not on the
size or molar mass of the molecules. Typical restrictions for accurate quantitative analysis
are that 1) the solutions are ideal 2) the solutions are dilute and 3) we only consider
nonelectrolytes.
vapor pressure lowering equation:
P   B PA
Note that change in P is directly proportional to the mole fraction of solute. This is simply
another way of looking at Raoult’s Law.
Since the vapor pressure is lowered with the presence of a solute and boiling is defined as
the temperature where the vapor pressure equals the atmospheric pressure one should
realize that solutions must have higher boiling points than the pure solvents. The equation
that is useful is:
Tb  Kb  m
Kb is the boiling point constant of the solvent and m is the solution molality.
Similarly, the freezing point is depressed. The equation is:
Tf  K f  m
Osmotic pressure, , the final colligative property is mathematically expressed:
  MRT 
n
RT
V
The derivation of this equation is to molarity since osmotic pressure measurements are
normally carried out at constant temperature. One calculation will allow you to see that very
dilute concentrations give rise to large osmotic pressures, especially when large proteins are
the solute. Thus, osmotic pressure measurements are a more sensitive method for
determining molar mass than the freezing point depression technique.
For colligative properties that are anomalous know the Van't Hoff factor, i. This factor hinges
on your ability to identify solutes that dissociate upon dissolving. Do you recognize ionic
compounds? Do you recognize strong acids – weak acids???